Lecture 12 Flashcards
bonds
localised interactions that hold a structure together
bond direction
dictates the shape that things adopt
bond strength
measure of how hard it is to break a bond, energy is needed to break a bond
matter
all matter is held together by bonds, changing the bond type changes its properties
solid
strong bonds, short range, complex networks in 3D that hold things in place
lots of bonds need to be broken to remove individual components
ordered
crystalline
disordered
amorphous
liquid
less bonds, weaker
bonds are held close together but no directionality or cooperation
molecules and atoms can slide, making bonds with molecules
gas
nothing bonds to anything else, free motionp
phase changes
the conversion from one phase to another
bonds are broken because energy (heat) is added
ionic solids
rigid 3D network
moving the ions can disrupt many long range bonds
hard, high melting, brittle
insulators, don’t conduct electricity
covalent networks
rigid 3D and 2D networks
hard, high melting, brittle
insulators
metals
metallic bonding
hard (strong bonds)
ductile (atoms slide but maintain bonding)
conduct electricity
intramolecular
strong covalent bonds inside molecules
intermolecular
van der waals and hydrogen bonding between molecules
covalent bonding
doesn’t take energy to disrupt the weaker bonds
soft, low melting
requires high temperature to start decomposing (break the covalent bond)
phase diagram
check notes
ionic bonding
electrostatic attraction due to opposite charges between electrons
ions interacting with each other
depend on distance, not direction
coulomb’s law
stronger for larger charges and smaller ions
ionic lattice
non-directional, larger, 3D arrangements of the ions, ionic lattice form
the packing maximises all of the attractive (-ve to +ve) contacts, as it minimises repulsive charges
the net extra bonding outweighs the extra repulsion
result of ionic solid
hard, strong, high melting solids
ions
large or small
mixtures of either or both
opposite charges is the only thing that matters
ions far apart
cations nucleus repels the anions nucleus and attracts its electrons
cations electrons repel the anions electrons and attract its nucleus
medium distances
the gap between them is more important (the atoms are tiny, % change or r is small)
the net positive and negative charges move the ions close together
ions close together
attraction doesn’t go away
the nuclei and some electrons are very close together (very strong repulsion)
the electrons that are further away from the nucleus (weaker attraction)
ionisation
energy required to remove an electron from an atom
electron gain enthalpy
energy required to add an electron to an ion
electron transfer example
check notes
positive energy
unfavourable
dangerous reaction that explodes because it gives out so much energy
electron transfer
occurs because electrons ‘fall’ from a higher energy orbital to a lower energy orbital
furthered stabilised by ionic bonds, which are very strong
anion always needs a cation
valence
charge an ion adopts depends on the position on the periodic table and the no of electrons used in bonding
valence orbitals
outer shell is the only one involved in ionic bonding
others are too high or low in energy
these are the valence shell
orbital energies
electron transfer is caused by orbital energy difference
the electron falls down that is the potential energy difference
electronegativity
measure of a degree to which an atom attracts an electron in a bond
increases across a period and decreases down a group
inverse relationship
as electronegativity increases, orbitals get lower in energy
number that describes the average energy of valence orbitals
vaan aarkel triangle
large x = low energy, contacted orbitals
small x = radially expanded, high energy orbitals
large enough electronegativity leads to an ionic bond.
